Calculate the heat evolved when 266 g of white phosphorous \( \left(\mathrm{P}_{4}\right) \) burns in air according to the equation: \( \mathrm{P}_{4}(s)+5 \mathrm{O}_{2}(\mathrm{~g}) \rightarrow \mathrm{P}_{4} \mathrm{O}_{10}(\mathrm{~s}) \quad \Delta H=-3013 \mathrm{~kJ} / \mathrm{mol} \)

When calculating the heat evolved from the combustion of white phosphorus, it's crucial to determine how many moles of \( P_4 \) are present in 266 g. The molar mass of \( P_4 \) is about 123.88 g/mol, so you can find the moles by dividing the mass by the molar mass. This gives you approximately 2.14 moles of \( P_4 \). Since the reaction releases \( -3013 \) kJ for each mole of \( P_4 \) consumed, you multiply the number of moles (2.14 moles) by the heat released per mole. So, the total heat evolved is roughly \( 2.14 \times -3013 \approx -6454 \) kJ. That's a lot of energy being released to the environment in flames!